1. Lab 9: Electrochemistry
Spring 2018
Lab 9: Electrochemistry

2. Goals Measure voltage of a galvanic cell
Determine how concentration affects voltage
Graphically determine value of Faraday constant

3. Oxidation-Reduction Reactions
A.k.a. Redox reactions
Involve electron transfer
Oxidation = Loss of electrons
Reduction = Gain of electrons
Balanced using half-reactions
Representations of oxidation and reduction

4. Half-Reactions Show how substances lose or gain electrons
Electrons written into reactions
Balance charge between reactants/products
Example: Fe(s) + Cd2+(aq) → Fe2+(aq) + Cd(s)
Fe(s) → Fe+2(aq) + 2e- (oxidation)
Cd2+(aq) + 2e- → Cd(s) (reduction)
0 charge =
0 charge

5. Balancing Redox Reactions
Balanced with half-reactions
Need same # of electrons in both so they cancel
Example:
Unbalanced Reaction: Fe(s) + Cl2(aq) → Fe3+(aq) + Cl-(aq)
Half-Reactions: Fe(s) → Fe3+(aq) + 3e-
Cl2(aq) + 2e- → 2Cl-(aq)
Balanced Half-Reactions: 2Fe(s) → 2Fe3+(aq) + 6e-
3Cl2(aq) + 6e- → 6Cl-(aq)
Balanced Reaction: 2Fe(s) + 3Cl2(aq) → 2Fe3+(aq) +6Cl-(aq)

6. Galvanic Cells
Produce voltage by separating half-reactions into half-cells
Necessary components:
Electrodes in matching solutions
Salt bridge
External circuit between electrodes to allow electron flow
Electron flow
Salt Bridge Cu Zn
Cu2+
Zn2+

7. Electrodes Solid surface for redox reaction Two types:
Anode – site of Oxidation (-)
Cathode – site of Reduction (+)
Current carried from anode to cathode
“The RED CAT gets FAT”

8. The Experiment The reaction: Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)
The set-up:
Half-cells in beral pipets
Agar plug salt bridge
Voltmeter attached to electrodes measures cell potential (Ecell)

9. Cell Potential Symbol: Ecell Units: Volts (V)
Related directly to Gibbs Free Energy:
n = # moles of transferred electrons (see half reactions)
F = Faraday’s constant (96,485 C/mol)
Potential in a galvanic cell is positive (spontaneous)

10. Standard Cell Potential
Symbol: E°cell
Units: Volts (V)
Voltage under standard conditions
1 M and 1 bar
Can be calculated for reaction using standard reduction potentials

11. Standard Reduction Potential (E°red)
Assigned to reduction half reactions
(e.g.) Cu2+(aq) + 2e- → Cu(s) E°red = V
See next slide for other values
Magnitude unaffected by balancing coefficients
I.e. Don’t multiply value!
Same number for reverse direction (oxidation)
Reversing direction leads to sign change.

12. More (+), more likely to be reduced
More (-), more likely to be oxidized

13. Calculating E°cell Combination of potentials for both half-cells:
E cell ° = E red ° − E ox °
Example: Fe(s) + Cl2(aq) → Fe3+(aq) + Cl-(aq)
E red °
Oxidation: 2Fe(s) → 2Fe3+(aq) + 6e V
Reduction: 3Cl2(aq) + 6e- → 6Cl-(aq) V
E cell ° = V – ( V) = V
Note: Eox is the same as Ered for a particular half-reaction (subtraction accounts for sign change associated with opposite direction).

14. Non-Standard Conditions
Potential (Ecell) affected by concentration
Concentrations used to calculate Ecell using the Nernst Equation:
R = gas constant (8.314 J/(K*mol)
T = temperature (K)
n = # moles of transferred electrons (see half reactions)
F = Faraday’s constant (96,485 C/mol)
Q = Reaction quotient that incorporates the concentrations

15. Reaction Quotient (Q) Same set-up as equilibrium constant expression:
Products as numerator raised to coefficient
Reactants as denominator raised to coefficient
Concentrations NOT at equilibrium
Can be initial concentrations
Example: N2O4(g) ⇌ 2NO2(g)

16. Determining F Graphically
Nernst equation can be graphed as a straight-line equation:
Slope used to calculate Faraday’s constant
y = b m * x
Experimentally—measured potential for the five trials
Natural log of ratio of solution concentrations for the five trials